A student investigates the effect of \(\text{pH}\) on the solubility of three hypothetical sparingly soluble salts: \(\text{MX}\), \(\text{MY}\), and \(\text{MZ}\). Each salt has a \(1:1\) cation-to-anion ratio and dissolves according to the generic equation:

\[ \text{MA(s)} \rightleftharpoons \text{M}^+\text{(aq)} + \text{A}^-\text{(aq)} \]

The student gathers the equilibrium data shown in the table below at \(298\text{ K}\).

Salt	\(K_{sp}\) at \(298\text{ K}\)	Anion	\(pK_b\) of Anion at \(298\text{ K}\)
\(\text{MX}\)	\(1.0 \times 10^{-10}\)	\(\text{X}^-\)	\(4.2\)
\(\text{MY}\)	\(1.0 \times 10^{-8}\)	\(\text{Y}^-\)	\(9.5\)
\(\text{MZ}\)	\(1.0 \times 10^{-6}\)	\(\text{Z}^-\)	\(16.2\)

Equal volumes of saturated solutions of each salt (each in contact with excess solid) are treated with a strong acid to lower the \(\text{pH}\) from \(7.0\) to \(2.0\). Which of the following correctly ranks the salts from greatest to least increase in molar solubility upon acidification, and provides the correct justification?